QCE Chemistry - Unit 3 - Oxidation and reduction

Redox Reactions | QCE Chemistry

Learn oxidation, reduction, oxidation states and redox equation reasoning for QCE Chemistry Unit 3.

Updated 2026-05-18 - 5 min read

Redox reactions are reactions where electrons are transferred from one species to another. The word "redox" combines reduction and oxidation because the two always happen together: one species loses electrons while another gains them.

Oxidation and reduction

A simple example is the reaction between zinc metal and copper(II) ions:

$ \mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)} $

Zinc starts as neutral metal and becomes $\mathrm{Zn^{2+}}$, so it has lost two electrons:

$ \mathrm{Zn(s)} \rightarrow \mathrm{Zn^{2+}(aq)} + 2e^- $

Copper(II) ions become neutral copper metal, so they have gained two electrons:

$ \mathrm{Cu^{2+}(aq)} + 2e^- \rightarrow \mathrm{Cu(s)} $

When the half-equations are added, the electrons cancel. This matters because electrons are transferred, not created or destroyed.

Redox is not a tiny isolated skill. QCE uses it to explain displacement reactions, combustion, corrosion, galvanic cells, electrolytic cells and practical observations. If a question mentions a metal reacting with ions, bubbling gas, plating on an electrode, rusting, or a colour change from one ion to another, it is often asking you to track electron transfer.

Oxidising and reducing agents

The names are easy to mix up, so slow down here:

• the oxidising agent causes oxidation and is itself reduced
• the reducing agent causes reduction and is itself oxidised

In the zinc-copper reaction, copper(II) ions are reduced because they gain electrons:

$ \mathrm{Cu^{2+}(aq)} + 2e^- \rightarrow \mathrm{Cu(s)} $

That means $\mathrm{Cu^{2+}}$ is the oxidising agent. Zinc loses electrons:

$ \mathrm{Zn(s)} \rightarrow \mathrm{Zn^{2+}(aq)} + 2e^- $

That means zinc is the reducing agent.

Oxidation states

Oxidation states are bookkeeping charges used to track electron movement, especially when electrons are not shown directly in the equation.

• If an element's oxidation state increases, it has been oxidised.
• If an element's oxidation state decreases, it has been reduced.
• Elements in their natural elemental form have oxidation state $0$, such as $\mathrm{Zn(s)}$, $\mathrm{Cu(s)}$, $\mathrm{O_2(g)}$ and $\mathrm{Cl_2(g)}$.
• A monatomic ion has an oxidation state equal to its charge, such as $\mathrm{Cu^{2+}} = +2$.
• Oxygen is usually $-2$, hydrogen is usually $+1$, and fluorine is always $-1$ in compounds.
• The total oxidation states in a neutral compound add to $0$. In a polyatomic ion, they add to the ion charge.

For $\mathrm{MnO_4^-}$, oxygen contributes $4 \times (-2) = -8$. The ion has total charge $-1$, so manganese must be $+7$.

Oxidants and reductants

The oxidant causes another species to be oxidised and is itself reduced. The reductant causes another species to be reduced and is itself oxidised.

In the zinc-copper example, $\mathrm{Cu^{2+}}$ is the oxidant because it accepts electrons from zinc. $\mathrm{Zn}$ is the reductant because it donates electrons to copper(II) ions.

This is the part students often find backwards: the oxidant is reduced, and the reductant is oxidised.

Balancing redox equations

In acidic solution, use this method for each half-equation:

1. Balance the key element being oxidised or reduced.
2. Balance oxygen by adding $\mathrm{H_2O(l)}$.
3. Balance hydrogen by adding $\mathrm{H^+(aq)}$.
4. Balance charge by adding $e^-$.
5. Multiply half-equations so the electrons cancel, then add and simplify.

This is often called the KOHES method: key element, oxygen, hydrogen, electrons, states. The order matters because adding water changes hydrogen, and adding hydrogen ions changes charge. Electrons should be the last balancing particle before you check states and combine.

For QCE, acidic conditions are the standard balancing context. That means you may add $\mathrm{H^+(aq)}$ and $\mathrm{H_2O(l)}$ while balancing. Do not invent hydroxide ions unless the question explicitly moves into alkaline conditions.

Using observations

Redox questions often give qualitative evidence instead of saying "this species was oxidised". Use the observations as clues:

• bubbles can mean a gas such as $\mathrm{H_2}$, $\mathrm{O_2}$ or $\mathrm{Cl_2}$ formed
• a metal coating can mean metal ions were reduced onto an electrode
• an electrode losing mass can mean metal atoms were oxidised into solution
• a colour fading or appearing can mean a coloured ion was consumed or produced

Good answers connect the observation to the electron transfer. For example, "the copper electrode gained mass" is useful, but "the copper electrode gained mass because $\mathrm{Cu^{2+}}$ ions were reduced to $\mathrm{Cu(s)}$ at the cathode" is much stronger.

Common mistake

Quick check

Sources

• QCAA Chemistry subject page
• QCAA Chemistry 2025 syllabus
• OpenStax Chemistry 2e: Review of Redox Chemistry
• OpenStax Chemistry 2e: Galvanic Cells